Batteries, overview

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Batteries, Overview From "Encyclopedia of Energy" © 2004, Elsevier Inc. 1. How Batteries Work 2. History 3. Types of Batteries and Their Characteristics 4. Applications 5. Environmental Issues 6. Future Outlook Glossary battery An assemblage of cells connected electrically in series and/or parallel to provide the desired voltage and current for a given application. cell An electrochemical cell composed of a negative electrode, a positive electrode, an electrolyte, and a container or housing. electrode An electronically conductive structure that provides for an electrochemical reaction through the change of oxidation state of a substance; may contain or support the reactant or act as the site for the electrochemical reaction. electrolyte A material that provides electrical conduction by the motion of ions only; may be a liquid, solid, solution, polymer, mixture, or pure substance. primary cell A cell that cannot be recharged and is discarded after it is discharged. secondary cell A rechargeable cell. Batteries are an important means of generating and storing electrical energy. They are sold at a rate of several billions of dollars per year worldwide. They can be found in nearly all motor vehicles (e.g., automobiles, ships, aircraft), all types of portable electronic equipment (e.g., cellular phones, computers, portable radios), buildings (as backup power supplies), cordless tools, flashlights, smoke alarms, heart pacemakers, biomedical instruments, wrist-watches, hearing aids, and the like. Batteries are so useful and ubiquitous, it is difficult to imagine how life would be without them. Batteries, strictly speaking, are composed of more than one electrochemical cell. The electrochemical cell is the basic unit from which batteries are built. A cell contains a negative electrode, a positive electrode, an electrolyte held between the electrodes, and a container or housing. Cells may be electrically connected to one another to form the assembly called a battery. In contrast to the alternating current available in our homes from the electric utility company, batteries deliver a direct current that always flows in one direction. There are a few different types of batteries: Primary batteries can be discharged only once and then are discarded; they cannot be recharged. Secondary batteries are rechargeable. Forcing current through the cells in the reverse direction can reverse the electrochemical reactions that occur during discharge. Both primary and secondary batteries can be categorized based on the type of electrolyte they use: aqueous, organic solvent, polymer, ceramic, molten salt, and so on. 1. HOW BATTERIES WORK The electrical energy produced by batteries is the result of spontaneous electrochemical reactions. The driving force for the reactions is the Gibbs free energy of the reaction, which can be calculated easily from data in tables of thermodynamic properties. The maximum voltage that a cell can produce is calculated from this simple relationship: where E is the cell voltage, ΔG is the Gibbs free energy change for the cell reaction (joules/mole), n is the number of electrons involved in the reaction (equivalents/mole), and F is the Faraday constant (96487 coulombs/equivalent). It is clear from Eq. (1) that a large negative value for the Gibbs free energy of the cell reaction is desirable if we wish to have a significant cell voltage. In practical terms, cell voltages of 1 to 2 V are achieved when using aqueous electrolytes, and cell voltages up to approximately 4 V are achieved when using nonaqueous electrolytes. When larger voltages are required, it is necessary to connect a number of cells electrically in series. For example, 12-volt automotive batteries are composed of six 2-volt cells connected in series. An important property of batteries is the amount of energy they can store per unit mass. This is the specific energy of the battery, usually expressed in units of watt-hours of energy per kilogram of battery mass (Wh/kg). The maximum value of the specific energy is that which can be obtained from a certain mass of reactant materials, assuming the case that any excess electrolyte and terminals have negligible mass. This is called the theoretical specific energy and is given by this expression: where the denominator is the summation of the molecular weights of the reactants. It can be seen from Eq. (2) that the theoretical specific energy is maximized by having a large negative value for ΔG and a small value for ΣMw. The value of ΔG can be made large by selecting for the negative electrode those reactant materials that give up electrons very readily. The elements with such properties are located on the left-hand side of the periodic chart of the elements. Correspondingly, the positive electrode reactant materials should readily accept electrons. Elements of this type are located on the right-hand side of the periodic chart. Those elements with a low equivalent weight are located toward the top of the periodic chart. These are useful guidelines for selecting electrode materials, and they help us to understand the wide interest in lithium-based batteries now used in portable electronics. Equations (1) and (2) give us a useful framework for representing the theoretical specific energy and the cell voltage for a wide range of batteries, as shown in Fig. 1. The individual points on the plot were calculated from the thermodynamic data. These points represent theoretical maximum values. The practical values of specific energy that are available in commercial cells are in the range of one-fifth to one-third of the theoretical values. As expected, the cells using high equivalent weight materials (e.g., lead, lead dioxide) have low specific energies, whereas those using low equivalent weight materials (e.g. lithium, sulfur) have high specific energies. The lines in Fig. 1 represent the cell voltages and simply represent the relationships given by Eqs. (1) and (2). FIGURE 1 Theoretical specific energy for various cells as a function of the equivalent weights of the reactants and the cell voltage. The details of the electrochemical reactions in cells vary, but the principles are common to all. During the discharge process, an electrochemical oxidation reaction takes place at the negative electrode. The negative electrode reactant (e.g., zinc) gives up electrons that flow into the electrical circuit where they do work. The negative electrode reactant is then in its oxidized form (e.g., ZnO). Simultaneously, the positive electrode reactant undergoes a reduction reaction, taking on the electrons that have passed through the electrical circuit from the negative electrode (e.g., MnO 2 is converted to MnOOH). If the electrical circuit is opened, the electrons cannot flow and the reactions stop. The electrode reactions discussed in the preceding can be written as follows: and The sum of these electrode reactions is the overall cell reaction: Notice that there is no net production or consumption of electrons. The electrode reactions balance exactly in terms of the electrons released by the negative electrode being taken on by the positive electrode. 2. HISTORY Volta's invention of the Volta pile in 1800 represents the beginning of the field of battery science and engineering. The pile consisted of alternating layers of zinc, electrolyte soaked into cardboard or leather, and silver. Following Volta's report, many investigators constructed electrochemical cells for producing and storing electrical energy. For the first time, relatively large currents at high voltages were available for significant periods of time. Various versions of the pile were widely used. Unfortunately, there was a lack of understanding of how the cells functioned, but as we now know, the zinc was electrochemically oxidized and the native oxide layer on the silver was reduced. The cells were “recharged” by disassembling them and exposing the silver electrodes to air, which reoxidized them. Inevitably, many other electrochemical cells and batteries were developed. John F. Daniell developed a two-fluid cell in 1836. The negative electrode was amalgamated zinc, and the positive electrode was copper. The arrangement of the cell is shown in Fig. 2. The copper electrodes were placed in (porous) porcelain jars, which were surrounded by cylindrical zinc electrodes and placed in a larger container. A copper sulfate solution was placed in the copper electrode's compartment, and sulfuric acid was put in the zinc electrode compartment. The electrode reactions were as follows: and FIGURE 2 A set of three Daniell cells connected in series. Sir William R. Grove, a lawyer and inventor of the fuel cell, developed a two-electrolyte cell related to the Daniell cell in 1839. Grove used fuming nitric acid at a platinum electrode (the positive) and zinc in sulfuric acid (the negative). Variants on this formulation were popular for a number of years. Of course, all of these cells were primary cells in that they could be discharged only once and then had to be reconstructed with fresh materials. Gaston Planté invented the first rechargeable battery in 1860. It was composed of lead sheet electrodes with a porous separator between them, spirally wound into a cylindrical configuration. The electrolyte was sulfuric acid. These cells displayed a voltage of 2.0 V, an attractive value. During the first charging cycles, the positive electrode became coated with a layer of PbO 2 . The charging operation was carried out using primary batteries—a laborious process. Because of the low cost and the ruggedness of these batteries, they remain in widespread use today, with various evolutionary design refinements. The electrode reactions during discharge of the lead-acid cell are as follows: and Notice that sulfuric acid is consumed during discharge and that the electrolyte becomes more dilute (and less dense). This forms the basis for determining the state of charge of the battery by measuring the specific gravity of the electrolyte. The theoretical specific energy for this cell is 175 Wh/kg, a rather low value compared with those of other cells. Waldemar Jungner spent much of his adult life experimenting with various electrode materials in alkaline electrolytes. He was particularly interested in alkaline electrolytes because there generally was no net consumption of the electrolyte in the cell reactions. This would allow for a minimum electrolyte content in the cell, minimizing its weight. Jungner experimented with many metals and metal oxides as electrode materials, including zinc, cadmium, iron, copper oxide, silver oxide, and manganese oxide. In parallel with Jungner's efforts in Sweden, Thomas Edison in the United States worked on similar ideas using alkaline electrolytes and many of the same electrode materials. Patents to these two inventors were issued at nearly the same time in 1901. During the period since the beginning of the 20th century, a wide variety of cells have been investigated, with many of them being developed into commercial products for a wide range of applications. Representative cells are discussed in the next section. 3. TYPES OF BATTERIES AND THEIR CHARACTERISTICS Batteries are usually categorized by their ability to be recharged or not, by the type of electrolyte used, and by the electrode type. Examples of various types of electrolytes used in batteries are presented in Table I. Primary (nonrechargeable) cells with aqueous electrolytes are usually the least expensive and have reasonably long storage lives. Historically, the most common primary cell is the zinc/manganese dioxide cell used for flashlights and portable electronics in sizes from a fraction of an amp-hour to a few amp-hours. Various modifications of this cell have been introduced, and now the most common version uses an alkaline (potassium hydroxide) electrolyte. The materials in this cell are relatively benign from an environmental point of view. The electrode reactions in an alkaline electrolyte may be represented as follows: and The cell potential is 1.5 V, and the theoretical specific energy is 312Wh/kg. Practical cells commonly yield 40 to 50Wh/kg. Another common primary cell is the zinc/air cell. It uses an aqueous potassium hydroxide electrolyte, a zinc negative electrode, and a porous catalyzed carbon electrode that reduces oxygen from the air. The overall electrode reactions are as follows: and TABLE I Electrolyte Types and Examples Aqueous electrolyte H 2 SO 4 , KOH Nonaqueous electrolyte Organic solvent Solid Li salt in ethylene carbonate-diethyl carbonate Crystalline Na 2 O x 11Al 2 O 3 Polymeric Li salt in polyethylene Molten salt (high temperature) LiCl-KCl The sum of these reactions gives this overall cell reaction: The cell voltage is 1.6 V, and the theoretical specific energy is 1200 Wh/kg. Practical specific energy values of 200 Wh/kg have been achieved for zinc/air cells. This cell has a very high energy per unit volume, 225Wh/L, which is important in many applications. The materials are inexpensive, so this cell is very competitive economically. This cell is quite acceptable from an environmental point of view. The zinc/mercuric oxide cell has the unique characteristic of a very stable and constant voltage of 1.35 V. In applications where this is important, this is the cell of choice. It uses an aqueous potassium hydroxide electrolyte and is usually used in small sizes. The electrode reactions are as follows: and The high equivalent weight of the mercury results in the low theoretical specific energy of 258 Wh/kg. Because mercury is toxic, there are significant environmental concerns with disposal or recycling. The zinc/silver oxide cell has a high specific energy of about 100 Wh/kg (theoretical specific energy =430 Wh/kg), and because of this, the cost of the silver is tolerated in applications that are not cost-sensitive. Potassium hydroxide is the electrolyte used here. The electrode reactions are as follows: and This gives the following overall cell reaction: The cell voltage ranges from 1.8 to 1.6 V during discharge. Because of the high density of the silver oxide electrode, this cell is quite compact, giving about 600Wh/L. Primary (nonrechargeable) cells with nonaqueous electrolytes have been under development since the 1960s, following the reports of Tobias and Harris on the use of propylene carbonate as an organic solvent for electrolytes to be used with alkali metal electrodes. It has long been recognized that lithium has very low equivalent weight and electronegativity (Table II). These properties make it very attractive for use as the negative electrode in a cell. Because lithium will react rapidly with water, it is necessary to use a completely anhydrous electrolyte. There are many organic solvents that can be prepared free of water. Unfortunately, organic solvents generally have low dielectric constants, making them poor solvents for the salts necessary to provide electrolytic conduction. In addition, organic solvents are thermodynamically unstable in contact with lithium, so solvents that react very slowly and form thin, conductive, protective films on lithium are selected. TABLE II Some Properties of Lithium and Zinc Lithium Zinc Equivalent weight (g/equiv.) 6.94 32.69 Reversible potential (V) -3.045 -0.763 Electronegativity 0.98 1.65 Density (g/cm 3 ) 0.53 7.14 Perhaps the most common lithium primary cell is that of Li/MnO 2 . It has the advantage of high voltage (compared with aqueous electrolyte cells), 3.05 V, and a high specific energy (~170 Wh/kg). The electrode reactions are as follows: and The electrolyte is usually LiClO 4 dissolved in a mixture of propylene carbonate and 1,2-dimethoxy-ethane. In general, the specific power that can be delivered by these cells is less than that available from aqueous electrolyte cells because the ionic conductivity of the organic electrolyte is much lower than that of aqueous electrolytes. An interesting cell is the one using fluorinated carbon as the positive electrode material. Because this material is poorly conducting, carbon and titanium current collection systems are used in the positive electrode. The electrode reaction is as follows: The electrolyte commonly is LiBF 4 dissolved in gamma butyrolactone. An unusual primary cell is that of lithium/thionyl chloride (SOCl2). The thionyl chloride is a liquid and can act as the solvent for the electrolyte salt (LiAlCl 4 ) and as the reactant at the positive electrode. This cell can function only because of the relatively stable thin protective film that forms on the lithium electrode, protecting it from rapid spontaneous reaction with the thionyl chloride. The cell reaction mechanism that is consistent with the observed products of reaction is the following: and where SO' is a radical intermediate that produces SO 2 and sulfur. This cell provides a high cell voltage of 3.6 V and a very high specific energy but can be unsafe under certain conditions. Specific energies of up to 700 Wh/kg have been achieved compared with the theoretical value of 1460 Wh/kg. Table III summarizes the characteristics of several representative primary cells. Primary cells with molten salt electrolytes are commonly used in military applications that require a burst of power for a short time (a few seconds to a few minutes). These batteries are built with an integral heating mechanism relying on a chemical reaction to provide the necessary heat to melt the electrolyte (at 400- 500°C). A typical cell uses lithium as the negative electrode and iron disulfide as the positive electrode. The following are example electrode reactions: and The molten salt electrolyte (a mixture of alkali metal chlorides) has a very high conductivity, allowing the cell to operate at very high specific power levels in excess of 1 kW/kg. Rechargeable cells with aqueous electrolytes have been available for more than 140 years, beginning with the Planté cell using lead and sulfuric acid as discussed previously. Modern Pb/PbO 2 cells and batteries have received the benefit of many incremental improvements in the design and optimization of the system. Current versions are very reliable and inexpensive compared with competitors. Depending on the design of the cell, lifetimes vary from a few years to more than 30 years. Specific energy values of up to approximately 40 Wh/kg are available. Alkaline electrolyte systems are available with a variety of electrodes. Perhaps the most common alkaline electrolyte cell is the Cd/NiOOH cell. It offers very long cycle life (up to thousands of charge-discharge cycles) and good specific power (hundreds of W/kg), albeit with a modest specific energy (35-55 Wh/kg). The cell voltage, 1.2 V, is lower than most and can be somewhat of a disadvantage. The electrode reactions are as follows: and Of course, the reverse of these reactions takes place on recharge. TABLE III Properties of Some Lithium Primary Cells Open circuit voltage (V) Practical specific energy (Wh/kg) Theoretical specific energy (Wh/kg) Li/MnO 2 3.0 280 750 Li/(CF) n 2.8 290 2100 Li/SOCl 2 3.6 450-700 1460 Li/FeS 2 1.7 220 760 A very robust rechargeable cell is the Edison cell, which uses iron as the negative electrode, nickel oxyhydroxide as the positive electrode, and an aqueous solution of 30w/o potassium hydroxide as the electrolyte. During the early years of the 20th century, Edison batteries were used in electric vehicles. They proved themselves to be very rugged and durable, although they did not have high performance. The iron electrode reaction is as follows: The positive electrode reaction is the same as for the Cd/NiOOH cell described earlier. The NiOOH electrode has proven itself to be generally the best positive electrode for use in alkaline electrolytes and has been paired with many different negative electrodes, including Cd, Fe, H 2 , MH, and Zn. Recently, the metal hydride/nickel oxyhydroxide cell (MH/NiOOH) has been a significant commercial success and has captured a large fraction of the market for rechargeable cells. It offers a sealed, maintenance-free system with no hazardous materials. The performance of the MH/NiOOH cell is very good, providing up to several hundred watts/kilogram peak power, up to about 85 Wh/kg, and several hundred cycles. The negative electrode operates according to the following reaction: The very high theoretical specific energy and low materials cost of the zinc/air cell make it attractive for consumer use. There have been many attempts to develop a rechargeable zinc/air cell, but with limited success due to the difficulties of producing a high-performance rechargeable air electrode. The electrode reactions during discharge for this cell are as follows: and During recent years, the cycle life of the air electrode has been improved to the point where more than 300 cycles are now feasible. Development efforts continue. In the meantime, various versions of a “mechanically rechargeable” zinc/air cell have been tested. These cells have provision for removing the discharged zinc and replacing it with fresh zinc. This approach avoids the difficulties of operating the air electrode in the recharge mode but creates the need for recycling the spent zinc to produce new zinc electrode material. The status of some rechargeable cells with aqueous electrolytes is shown in Table IV. Rechargeable cells with nonaqueous electrolytes have been under development for many years, although they have been available on the consumer market for only a decade or so. All of the types of nonaqueous electrolytes shown in Table I have been used in a variety of systems. Organic solvent-based electrolytes are the most common of the nonaqueous electrolytes and are found in most of the rechargeable lithium cells available today. A typical electrolyte consists of a mixture of ethylene carbonate and ethyl-methyl carbonate, with a lithium salt such as LiPF 6 dissolved in it. Various other solvents, including propylene carbonate, dimethyl carbonate, and diethyl carbonate, have been used. Other lithium salts that have been used include lithium perchlorate, lithium hexafluoro arsenate, and lithium tetrafluoro borate. All of these combinations of solvents and salts yield electrolytes that have much lower conductivities than do typical aqueous electrolytes. As a result, the electrodes and electrode spacing in these lithium cells are made very thin to minimize the cell resistance and maximize the power capability. Lithium is difficult to deposit as a smooth compact layer in these organic electrolytes, so a host material, typically carbon, is provided to take up the Li on recharge and deliver it as Li ions to the electrolyte during discharge. Many types of carbon, both graphitic and nongraphitic, have been used as the lithium host material. In addition, a variety of intermetallic compounds and metals have been used as Li host materials. All of the commercial Li rechargeable cells today use a carbon host material as the negative electrode. TABLE IV Rechargeable Aqueous Battery Status System Cell voltage (V) Theoretical specific energy (Wh/kg) Specific energy (Wh/kg) Specific power (W/kg) Cycle life Cost (dollars/kWh) Pb/PbO 2 2.10 175 30-45 50-100 >700 60-125 Cd/NiOOH 1.20 209 35-55 400 2000 >300 Fc/NiOOH 1.30 267 40-62 70-150 500- 2000 >100 H 2 /NiOOH 1.30 380 60 160 1000- 2000 >400 Note. est., estimate. [...]... http://search.credoreference.com/content/entry/estenergy/batteries _overview/ 0 APA CAIRNS, E.(2004) Batteries, overview In Encyclopedia of energy Retrieved from http://search.credoreference.com/content/entry/estenergy/batteries _overview/ 0 MLA CAIRNS, ELTON J "Batteries, Overview. " Encyclopedia of Energy Oxford: Elsevier Science & Technology, 2004 Credo Reference Web 14 May 2014 Chicago CAIRNS, ELTON J "Batteries, Overview. " In Encyclopedia... lithium/fluorine cell could be developed, this would lead to an even higher specific energy SEE ALSO THE FOLLOWING ARTICLES Batteries, Transportation Applications • Electrical Energy and Power • Electric Motors • Storage of Energy, Overview Further Reading Linden, D.; Reddy, T B (eds.) (2002) “Handbook of Batteries, 3rd ed McGraw-Hill New York Socolow, R.; Thomas, V (1997) The industrial ecology of lead and electric... ELTON J "Batteries, Overview. " In Encyclopedia of Energy Oxford: Elsevier Science & Technology, 2004 http://search.credoreference.com/content/entry/estenergy/batteries _overview/ 0 (accessed May 14, 2014.) Harvard CAIRNS, E 2004 'Batteries, overview' in Encyclopedia of energy, Elsevier Science & Technology, Oxford, United Kingdom Accessed: 14 May 2014, from Credo Reference ... related functions Golf carts, utility vehicles, and forklift trucks represent a variety of propulsion-power applications There are many government and military applications with stringent requirements for batteries, including communications radios and phones, portable lasers for range finding, field computers, night vision equipment, sensors, starting for all types of vehicles, power sources for fuses, . http://search.credoreference.com/content/entry/estenergy/batteries _overview/ 0 APA CAIRNS, E.(2004). Batteries, overview. In Encyclopedia of energy. Retrieved from http://search.credoreference.com/content/entry/estenergy/batteries _overview/ 0 MLA CAIRNS,. J. " ;Batteries, Overview. " Encyclopedia of Energy. Oxford: Elsevier Science & Technology, 2004. Credo Reference. Web. 14 May 2014. Chicago CAIRNS, ELTON J. " ;Batteries, Overview. ". ARTICLES Batteries, Transportation Applications • Electrical Energy and Power • Electric Motors • Storage of Energy, Overview Further Reading Linden, D.; Reddy, T. B. (eds.). (2002). “Handbook of Batteries,